Archive for the 'Kass' Category

Kass and coworkers looked at a series of substituted phenols to tease out ways to produce stronger acids in non-polar media.1 First they established a linear relationship between the vibrational frequency shifts of the hydroxyl group in going from CCl4 as solvent to CCl4 doped with 1% acetonitrile with the experimental pKa in DMSO. They also showed a strong relationship between this vibrational frequency shift and gas phase acidity (both experimental and computed deprotonation energies).

A key recognition was that a charged substituent (like say ammonium) has a much larger effect on the gas-phase (and non-polar solvent) acidity than on the acidity in a polar solvent, like DMSO. This can be attributed to the lack of a medium able to stable charge build-up in non-polar solvent or in the gas phase. This led them to 1, for which B3LYP/6-31+G(d,p) computations of the analogous dipentyl derivative 2 (see Figure 1) indicated a deprotonation free energy of 261.4 kcal mol-1, nearly 60 kcal mol-1 smaller than any other substituted phenol they previously examined. Subsequent measurement of the OH vibrational frequency shift showed the largest shift, indicating that 1 is extremely acidic in non-polar solvent.

Further computational exploration led to 3 (see Figure 1), for which computations predicted an even smaller deprotonation energy of 231.1 kcal mol-1. Preparation of 4 and experimental observation of its vibrational frequency shift revealed an even larger shift than for 1, making 4 extraordinarily acidic.

2

Conjugate base of 2

3

Conjugate base of 3

Figure 1. B3LYP/6-31+G(d,p) optimized geometries of 2 and 3 and their conjugate bases.

The concept of antiaromaticity is an outgrowth of the well-entrenched notion or aromaticity. While 4n+2 π-electron systems are aromatic, 4n π-electron systems should be antiaromatic. That should mean that antiaromatic systems are unstable. The cyclopropenyl anion 1a has 4 π-electrons and should be antiaromatic. Kass has provided computational results that strongly indicate it is not antiaromatic!1

Let’s first look at the 3-cyclopropenyl cation 1c. Kass has computed (at both G3 and W1) the hydride affinity of 1c-4c. The hydride affinities of the latter three compounds plotted against the C=C-C+ angle is linear. The hydride affinity of 1c however falls way below the line, indicative of 1c being very stable – it is aromatic having just 2 π-electrons.

A similar plot of the deprotonation enthalpies leading to 1a-4d vs. C=C-C– angle is linear including all four compounds. If 1a where antiaromatic, one would anticipate that the deprotonation energy to form 1a would be much greater than expected simply from the effect of the smaller angle. Kass suggests that this indicates that 1a is not antiaromatic, but just a regular run-of-the-mill (very) reactive anion.

A hint at what’s going on is provided by the geometry of the lowest energy structure of 1a, shown in Figure 1. The molecule is non-planar, having Cs symmetry. A truly antiaromatic structure should be planar, really of D3h symmetry. The distortion from this symmetry reduces the antiaromatic character, in the same way that cyclobutadiene is not a perfect square and that cyclooctatraene is tub-shaped and not planar. So perhaps it is more fair to say that 1a has a distorted structure to avoid antiaromaticity, and that the idealized D3h structure, does not exist because of its antiaromatic character.

The difference in the acidity of t-butanol and 2, some 30 kcal mol-1, reflects the stability afforded by three intramolecular hydrogen bonds to the oxyanion. In going from 2cb to 1cb, each of the hydroxyl groups that donate to the oxyanion act as the acceptor of a hydrogen bond from the more removed hydroxyl groups. There is in effect a first and second layer of hydrogen bond network in 1cb. These secondary hydrogen bonds lead to further stabilization of the anion, as reflected in the diminished DPE of 1 over 2: 320.2 vs. 335.0 kcal mol-1. Note that this secondary layer does not stabilize the anion to the same degree as the primary layer, but nonetheless its effect is large and quite striking.

1cb

Figure 1. M06-2x/maug-cc-pVT(+d)Z optimized structure of 1cb.

Even in solution these more remote hydrogen bonds can stabilize the anion. So, using the CPCM approach and modeling DMSO, 2 is predicted have a pKa that is 15 units below that of t-butanol, and 1 is predicted to be 3 pKa units more acidic than 2. Experiments verify this prediction with the pKas of 16.1 for 2 and 11.4 for 1.

I have had a long-standing interest in organolithium compounds, dating back to my graduate student days. Thus, I was excited to read Kass and Radom’s latest work on the computational and experimental evaluation of the acidity of lithium acetylide LiCCH.1

The gas phase experimental acidity is accomplished by preparing the conjugate base of lithium acetylide through a procedure of collision-induced dissociation with loss of CO2, as in Scheme 1. By reacting this anion with a variety of different acids, they were able to bracket the acidity and determine that ΔHacid is 391.0 ± 1.3 kcal mol-1. This is about 13 kcal mol-1less acidic than acetylene itself. The reduction is acidity understandable in terms of the C-Li being essentially ionic, and thereby loss of the proton builds up negative charge on a carbon adjacent to a carbon that already has a great deal of negative charge.

Scheme 1

Computations support this enthalpy for deprotonation. The G3, G4 and W1 values for the enthalpy deprotonation of lithium acetylide are 389.1, 388.9, and 390.4 kcal mol-1, respectively. It should also be noted that the conjugate base of lithium acetylide posses a non-classical bridging geometry 1, which is well-known for organolithium species.2

The α-proton of ketones and aldehydes are acidic, thanks to delocalization of the resulting anion. However, α-protons at a bridgehead position are much less acidic – the resulting anion is not delocalized as the enolate would be an anti-Bredt alkene. So, what about more remote protons from the carbonyl – would they exhibit enhanced acidity due to inductive or field effects?

Kass has examined the deprotonation of 2-adamantone 1 via experiment and computation.1 The relative energies of the five different anions are listed in Table 1. Previous H/D exchange experiments indicate that the relative reactivity is βax > βeq > α, and this is well reproduced by computations.2

Table 1. Relative energies (kcal mol-1) of the enolates of 1.

compound

M06-2x/aug-cc-pVDZ

G3

α

4.27

5.60

βax

0.0

0.0

βeq

4.46

γ

2.28

3.40

δ

6.17

7.55

2

-1.58

0.56

Kass’ bracketing experiments indicate the enthalpy for deptrotonation of 2-adamantone is 394.7 ± 1.4 kcal mol-1. This is in nice accord with the computational results for loss of the βax proton: 393.8 (M06-2x/aug-cc-pVDZ) and 396.8 kcla mol-1 (G3). One interesting computational result is a competive cyclic structure 2, whose stability is similar to that to the βax ion at M06-2x and is the optimized structure produced at MP2/6-31G(d) when searching for the βeq enolate.

So, to answer our question, protons remote from a carbonyl are more acidic than alkane
analogues, but much less acidic than typical α-protons of ketones.

Molecular structures can differ depending on phase, particularly between the gas and solution phase. Kass has looked at the protonation of 4-aminobenzoic acid. In water, the amino is its most basic site, but what is it in the gas phase? The computed relative energies of the protonation sites are listed in Table 1. If one corrects the B3LYP values for their errors in predicting the proton affinity of aniline and benzoic acid, the carbonyl oxygen is predicted to be the most basic site by 5.0 kcal mol-1, in nice accord with the G3 prediction of 4.1 kcal mol-1. Clearly, the structure depends on the medium.

Electrospray of 4-aminobenzoic acid from 3:1 methanol/water and 1:1 acetonitrile/water solutions gave different CID spectra. H/D exchange confirmed that electrospray from the emthanol/water solution gave the oxygen protonated species while that from the acetonitrile/water solution gave the ammonium species.

Following on their prediction that the thiol of cysteine1 is more acidic than the carboxylic acid group (see this post), Kass has examined the acidity of tyrosine 1.2 Which is more acidic: the hydroxyl (leading to the phenoxide 2) or the carboxyl (leading to the carboxylate 3) proton?

1

2

3

Kass optimized the structures of tyrosine and its two possible conjugate bases at B3LYP/aug-cc-pVDZ, shown in Figure 1, and also computed their energies at G3B3. 2 is predicted to be 0.2 kcal mol-1 lower in energy than 3 at B3LYP and slightly more stable at G3B3 (0.5 kcal mol-1). However, both computational methods underestimate the acidity of acetic acid more than that of phenol. When the deprotonation energies are corrected for this error, the phenolic proton is predicted to be 0.4 kcal mol-1 more acidic than the carboxylate proton at B3LYP and 0.9 kcal mol-1 more acidic at G3B3.

1

2

3

Figure 1. B3LYP/aug-cc-pVDZ optimized structures of tyrosine 1 and its two conjugate bases 2 and 3.2

Gas phase experiments indicate that deprotonation of tyrosine leads to a 70:30 mixture of the phenoxide to carboxylate anions. The computations are in nice agreement with this experiment. (A Boltzmann weighting of the computed lowest energy conformers makes only a small difference to the distribution relative to using simply the single lowest energy conformer.) This demonstrates once again the important role of solvent, since only the carboxylate anion is seen in aqueous solution.

Kass has once again uncovered a simple system that challenges our notions of basic chemical concepts. It is a well accepted notion that the most acidic proton of all of the amino acids is the carboxylic acid one. However, acidities are strongly influenced by the solvent, and the absence of solvent in the gas phase can dramatically alter things.

Kass and co-workers examined the gas-phase acidity of cysteine with computational and
experimental techniques.1 The lowest energy conformer of cysteine is 1a, characterized by having three intramolecular hydrogen bonds (Figure 1). The next lowest conformer, 1b, has only two intramolecular hydrogen bonds and is 1.5 kcal mol-1 higher in energy at G3B3.

They optimized a number of different configurations of the conjugate base of cysteine: two conformers from the loss of the carboxylate proton (2a and 2b), two conformers from the loss of the thiol proton (2c and 2d), and one conformer from the loss of the thiol proton of the zwitterion (2e). These structures are shown in Figure 2 along with their relative energies. All of these structures possess two intramolecular hydrogen bonds.

The gas phase acidity of carboxylic acids is greater than thiols; the deprotonation energy of propanoic acid (CH3CH2CO2H) is 347.7 kcal mol-1 at G3B3 (347.2 expt.2), about 6 kcal mol-1 less than that of ethanethiol (CH2CH2SH: 355.0 at G3B3 and 354.2 expt.2). However, the computations indicate that 2c is the lowest energy structure of deprotonated cysteine, and 2c comes about by loss of the thiol proton! Te lowest energy cysteine conjugate base from loss of the carboxylate proton is 1a, which is 3.1 kcal mol-1 higher in energy. Apparently, the hydrogen bonding network in 2c is quite favorable, able to make up for the inherent favorability of a carboxylate over a thiolate anion.

The G3B3 computed deprotonation energy of cysteine is 333.3 kcal mol-1 (for removal of the thiol proton). Kass determined the deprotonation energy of cysteine using a kinetic and a thermodynamic method. The kinetic method gives a value of 332.9 ± 3.3 kcal mol-1­, while the thermodynamic method gives 334.4 ± 3.3 kcal mol-1­. These are in fine agreement with the computed value.

This study ably demonstrates the dramatic role that solvent can play in determining molecular properties. Kass titled the article “Are carboxyl groups the most acidic sites in amino acids?” and answers with “no” – in the gas phase the thiol group is more acidic. He ends the article with an indication that the alcohol of tyrosine may be competitive in acidity with its carboxylic group, too.

InChIs

In Chapter 2.2, we suggest that the experimental deprotonation energy (DPE) of cyclohexane is in doubt. G2MP2 predicts the DPE of cyclohexane is 414.5 kcal mol-1, a figure significantly higher than the experimental1 value of 404 kcal mol-1. Given that the deviation between the G2MP2 computed DPE and experiment is about 2 kcal mol-1, we suggest that cyclohexane should be re-examined.

In a recent JACS article,2 Kass calls into question the experimental bond dissociation energies (BDE) of the small cycloalkanes. With his experimental determination of the BDE of both the vinyl and allylic positions of cyclobutene, Kass can compare experimental and computed BDEs for a range of hydrocarbon environments, as listed in Table 1. The two composite methods G3 and W1 provide excellent BDE values for the small alkanes, one acyclic alkene, and the small cyclic alkenes. These composite methods appear to accurately predict BDEs of hydrocarbons.

However, the small cyclic alkanes are dramatic outliers. The well-accepted experimental BDEs of cyclopropane, cyclobutane, and cyclohexane are 3-5 kcal mol-1 lower than those predicted by the composite methods. Given the strong performance of the computational methods, and the difficulties associated with experimental determinations of BDEs, Kass suggests that the BDEs of these cycloalkanes are in error. Further experiments are deserved.